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Bureau of Mines Information Circular/1981 




Selection of Lixiviants for In Situ 
Uranium Leaching 

By Daryl R. Tweeton and Kent A. Peterson 




UNITED STATES DEPARTMENT OF THE INTERIOR 



Information Circular 8851 

Selection of Lixiviants for In Situ 
Uranium Leaching 

By Daryl R. Tweeton and Kent A. Peterson 




UNITED STATES DEPARTMENT OF THE INTERIOR 
James G. Watt, Secretary 

BUREAU OF MINES 




As the Nation's principal conservation agency, the Department of the Interior 
has responsibility for most of oiir nationally owned public lands and natural 
resources. This includes fostering the wisest use of our land and water re- 
sources, protecting our fish and wildlife, preserving the environmental and 
cultural values of our national parks and historical places, and providing for 
the enjoyment of life through outdoor recreation. The Department assesses 
our energy and mineral resources and works to assure that their development is 
in the best interests of all our people. The Department also has a major re- 
sponsibility for American Indian reservation communities and for people who 
live in Island Territories under U.S. administration. 




. Ua 



This publication has been cataloged as follows: 



Tweeton, Daryl R 

Selection of lixiviants for in situ uranium leaching. 

(Information circular / Bureau of Mines ; 8851) 
Bibliography: p. 23*25. 

1. Uranium mines and mining. 2. In situ processing (Mining). 3. 
Leaching. 4. Solvents. I. Peterson, Kent A. 11. Title. III. Information 
circular (United States. Bureau of Mines) ; 8851. 



TN295.U4 [TN490.U7] 622s [622'.34932] 81-6103 AACR2 



For sale by the Superintendent of Documents, U.S. Government Printing Office 
Washington. D.C. 20402 



CONTENTS 

Page 

Abstract 1 

Introduction 1 

Acknow! edgments 2 

Aval 1 abl e 1 i xi vi ants 2 

Alkaline (carbonate-bicarbonate) lixiviants 6 

Ammonium carbonate-bicarbonate 6 

Sodium carbonate-bicarbonate 7 

Potassium carbonate-bicarbonate 9 

Carbon dioxide 9 

Acid lixiviants 11 

Sul furic acid 11 

Nitric and hydrochloric acids 12 

Oxidizers 13 

Oxygen 13 

Hydrogen peroxide 14 

Chi orates 14 

Methods of testing lixiviants 15 

Laboratory tests 15 

Batch 1 eachi ng 15 

Col umn 1 eachi ng 16 

Pilot field tests 17 

Types of tests 17 

Problems and solutions 18 

Geochemical model s 19 

Summary 22 

References 23 

Appendi x 26 

TABLES 

1. Costs of chemicals used for making lixiviants 3 

2. Summary of alkaline lixiviant costs, advantages, and disadvantages 4 

3. Summary of acid lixiviant costs, advantages, and disadvantages 5 

4. Summary of oxidizer costs, advantages, and disadvantages 5 

A-1. Initial (analytical) data describing solution 27 

A-2. Summary of differences between analytical and computed solution 28 

description 

A-3. Computed distribution of species in solution 29 

A-4. Degree of saturation of mineral phases 33 



SELECTION OF LIXIVIANTS FOR IN SITU URANIUM LEACHING 

by 

Daryl Ri Tweeton ^ and Kent A. Peterson2 



ABSTRACT 



This Bureau of Mines publication 
provides information to assist in select- 
ing a lixiviant (leach solution) for in 
situ uranium leaching. The cost, advan- 
tages, and disadvantages of lixiviants 
currently used and proposed are pre- 
sented. Laboratory and field tests are 
described, and applications of geochemi- 
cal models are discussed. 

Environmental, economic, and techni- 
cal factors should all be considered. 
Satisfying environmental regulations on 
restoring groundwater quality is becoming 
an overriding factor, favoring sodium 
bicarbonate or dissolved carbon dioxide 
over ammonium carbonate. The cheapest 



lixiviant is dissolved carbon dioxide, 
but it is not effective in all deposits. 
Technical factors include clay swelling 
by sodium, acid consumption by calcite, 
and the low solubility of oxygen in shal- 
low deposits. 

Laboratory leaching tests can pro- 
vide useful data. However, they can be 
misleading if, for example, the ore is 
allowed to oxidize before testing or if 
distilled water instead of formation 
water is used for making solutions for 
permeability tests. Geochemical models 
presently are more useful for indicating 
trends in solubility than in reliably 
predicting concentrations. 



INTRODUCTION 



The Bureau of Mines, as part of its 
mission to help assure that the Nation 
has an adequate supply of minerals, con- 
ducts research on alternative recovery 
techniques such as in situ leaching. An 
important consideration in any in situ 
leaching operation is the choice of lix- 
iviant (the leach solution). This publi- 
cation is intended to assist companies 
planning an in situ uranium leaching 
operation to select a lixiviant. General 
knowledge of in situ leaching is assumed; 

^Research physicist. Twin Cities Research 
Center, Bureau of Mines, Twin Cities, 
Minn. 
Geologist, Twin Cities Research Center, 
Bureau of Mines, Twin Cities, Minn, 



readers wanting a basic introductory pub- 
lication should consult other literature 
(12-15^, 30). 3 

Selection of the lixiviant is of 
critical importance to the success of an 
in situ leaching operation. The lixivi- 
ant affects not only the recovery of ura- 
nium and the cost of chemicals, but also 
the difficulty of meeting environmental 
regulations concerning restoration of 
groundwater quality after leaching. 
Thus, the choice of lixiviant is not 



'Underlined numbers in parentheses refer 
to items in the list of references 
preceding the appendix. 



simple, as technical , 
latory factors should 



economic, and regu- 
all be considered. 



information 
improve this 



or made 
report. 



suggestions to help 



No manuscript specifically on lixiv- 
iant selection has previously been made 
available to the public. Much research 
has been done by companies, but the 
results have usually been considered pro- 
prietary. However, useful literature is 
available on topics that are important 
parts of the lixiviant selection process. 
The chemistry of conventional uranium 
milling is thoroughly discussed by 
Merritt (19^). The similarities in chem- 
istry between milling and in situ leach- 
ing make this a yery useful reference. 
However, differences such as the gener- 
ally higher concentrations and consump- 
tions of chemicals and higher recovery of 
uranium in milling should be kept in mind 
when considering in situ leaching. (The 
section on in situ leaching does not 
represent the current state of the art 
because this reference was published in 
1971.) Extensive column leaching studies 
were performed for the Bureau of Mines by 
Westinghouse Electric Corp. (1£). These 
reports compare results using various 
lixiviants on two sandstone ores and pro- 
vide recommendations concerning methods 
of conducting column leaching tests in 
general. The lixiviants were usually 
made with distilled water instead of com- 
patible formation water, so permeability 
results should be interpreted with cau- 
tion. The influence of various lixivi- 
ants on the difficulty of restoring the 
groundwater quality after leaching is an 
important factor and is discussed in sev- 
eral publications (13, 30). Well drill- 
ing and completion TT, 75", 33-34) affect 
lixiviant selection i n'^Trec tTy , as good 
injection and production wells are vital 
to the success of field tests. 

ACKNOWLEDGMENTS 

The authors wish to thank Rocky 
Mountain Energy Co. for providing access 
to data concerning sulfuric acid leaching 
and to FMC Corp. for supplying informa- 
tion about hydrogen peroxide and sodium 
sesquicarbonate. The authors are also 
grateful to the many individuals, too 
many to list separately, who contributed 



AVAILABLE LIXIVIANTS 

Lixiviants that have been used for 
in situ uranium leaching include solu- 
tions of ammonium carbonate-bicarbonate, 
sodium carbonate-bicarbonate, carbon 
dioxide, and sulfuric acid. Potassium 
carbonate-bicarbonate is technically 
attractive but has been considered too 
expensive. Hydrochloric and nitric acids 
have been proposed for leaching carbona- 
ceous ore. The carbonate-bicarbonate and 
sulfuric acid lixiviants contain an anion 
that will form a soluble complex with 
uranium in its +6 charge state. The 
cation does not directly affect the solu- 
bility of the uranium but is important 
because of its effect on permeability, 
cost, and groundwater quality 
restoration. 

An oxidizer is required to convert 
unoxidized uranium from its insoluble 
+4 charge state to its soluble +6 charge 
state. Oxidizers that have been used 
include oxygen, hydrogen peroxide, and 
sodium chlorate. 

The costs, not including delivery, 
of chemicals used for making lixiviants 
are listed in table 1 in the forms com- 
monly used in dollars per kilogram and 
per ton. Costs were obtained from dis- 
cussions with suppliers and leaching com- 
panies and from published prices U) in 
late 1980. These units are not the most 
useful for comparing costs because they 
do not directly compare the cost of pro- 
viding the significant component. For 
example, 1 kg of potassium carbonate pro- 
vides less carbonate that 1 kg of ammo- 
nium carbonate. 



To facilitate compari 
tables 2, 3, and 4 present 
kilogram-mole and a pound- 
line lixiviants, acid li 
oxidizers, respectively, 
the costs of bicarbonate 
lixiviants separately to pe 
ing the cost of lixiviant 
proportion of the two. 



ng the costs, 

the cost of a 

mole of alka- 

xiviants, and 

Table 2 lists 

and carbonate 

rmit calculat- 

containing any 

Costs of 



lixiviants of typical concentrations are 
also given in terms of cents per liter 
and cents per gallon in the discussion of 
each lixiviant that follows the summary 
tables. These lixiviant costs were cal- 
culated using the costs of chemicals in 



table 1, so changes in those chemical 
prices will cause proportionate changes 
in the corresponding lixiviant costs. 
The cost of $80/ton of carbon dioxide 
used in formulating table 2 is typical, 
but it can vary a great deal. 



TABLE 1. - Costs of chemicals used for making lixiviants 



Cost, 100-pct basTT 



Chemical 



Form 



Per ton 



Per kg 



CO. 



NH: 



Compressed and cooled liquefied gas' 
Ammonia fertil izer 



NaOH.., 
Na2C03, 
NaHCO,, 



NaHC03 
Na2C03 2H2O. 



Caustic soda, liquid, 50 pet at $250/ton. 

Soda ash 

Fl akes or powder 

Sodium sesquicarbonate granules , 



KOH. 



K2CO3, 
KHCO,, 



H2SO, 
HNO3, 
HCl., 



Caustic potash, liquid, 45 pet KOH 

at $11/100 lb. 
Liquid, 47 pet K2CO3 at $11.75/100 lb. 
Granulated, technical grade 



Liquid, concentrated virgin.., 

Liquid, 58.5-68 pet HNO3 , 

Liquid, 37 pet HCl at $70/ton, 



Liquid, at $0.40-$0. 60/100 cu ft of O2 gas 
at 1 atm and 25** C (1 lb = 12.08 cu ft). 



H2O2.. 
NaClO- 



Liquid, 50 pet H2O2 
Powder or flakes.... 



at $0.29/lb. 



$60-$ 200 
200 

500 
90 

240 
96 



490 

500 
280 

80 
120 
190 

297 
5145 

1160 
400 



$0.07-0.22 

.22 

.55 
.10 
.26 
.11 



.54 

.55 
.31 

.09 
.13 
.21 

2.11 
3.16 

1.28 
.44 



'Depending on annual use. 

2Texas. 

^Wyoming. 



See text. 



TABLE 2. - Summary of alkaline lixiviant costs, advantages, and disadvantages 



Cost 



Lixiviant 



Ammonium lixiviants: 

NH4HCO3 , 

(NH4)2C03 



Per kg-mole 



Per lb-mole 



PPm U3O8 
paying^ for 
3 g/1 anion 



Advantages and 
disadvantages 



Little effect on 
permeability. 
Difficult to 
meet restoration 
requirements. 



Relatively easy 
to meet resto- 
ration require- 
ments. Can 
reduce perme- 
ability. 



Sodium lixiviants: 
From soda ash: 

NaHCOj 

Na2C03 

From caustic soda: 

NaHC03 

Na2C03 

From sodium 
sesquicarbonate: 

NaHC03 

Na2C03 

Potassium lixiviants: 
From caustic potash: 

KHCO3 

K2CO3 

From granules: 

KHCO3 

From 47% K2CO3 solution 

K2CO3 



CO^ 



$7.63 
11.38 



7.19 
10.52 

25.93 
47.98 



8.67 
17.60 



33.56 
63.26 

30.90 

76.18 



23.88 



$3.46 
5.16 



3.27 
4.77 

11.76 
21.76 



3.93 
7.98 



15.22 
28.69 

14.00 

34.55 



21.76 



5.7 
8.6 



5.4 
8.0 



19 
36 



6.5 
13 



25 
48 

23 

58 



2»32.9 



Little effect on 
permeability, 
should be rela- 
tively easy to 
meet restoration 
requirements. 
Expensive unless 
preceded by 
chloride preflush. 

Cheap, little 
effect on perme- 
ability, easy to 
meet restoration 
requirements. ■ 
Not effective on 
all deposits. 



^Assuming $66/kg for U3O8 and neglecting recycling. 
2Assuming $80/ton. 
^For 3 g/1 HCO3. 



TABLE 3. - Summary of acid lixiviant costs, advantages, and disadvantages 





Cost 










ppm U3O8 


. Advantages and 


Lixiviant 


Per kg-mole 


Per lb-mole 


paying! for 


disadvantages 




(per kg-equiv wt) 


(per Ib-equiv wt) 


0.051 M solution2 
(0.102 ¥ solution)2 




hi ooU^ .... 


$8.65 


$3.92 


6.7 


Very effective in 




(4.33) 


(1.96) 


(6.7) 


amenable depos- 
its, restoration 
easier than with 
(NH4)2C03. Not 
usable in depos- 
its with much 
CaC03, not selec- 
tive for U. 


HNO3 


8.33 


3.78 


6.4 


Claimed to be 




(8.33) 


(3.78) 


(13) 


effective for 
carbonaceous 
deposits. Not 
selective for U, 
dissolves Ra, 
requires cationic 
IX resin, diffi- 
cult restoration. 


HCl 


7.61 


3.45 


5.9 


Claimed to be 




(7.61) 


(3.45) 


(12) 


effective for 
clayey deposits. 
Not selective for 
U, dissolves Ra, 
requires cationic 
IX resin. 



^Assuming $66/kg for U3O8 and neglecting recycling. 
^Equivalent to 5 g/1 H2SO4. 

TABLE 4. - Summary of oxidizer costs, advantages, and disadvantages 





Cost 




Oxidizer 


Per kg-mole 


Per lb-mole 


ppm U3O8 
paying! for 
0.3 g/1 O2 


Advantages and 
disadvantages 


O2: 










Texas 

Wyoming... 


$3.40 
5.10 


$1.55 
2.33 


0.48 
.72 


Cheap. Must be injected down- 
hole, can cause gas blockage 
near injection wells. 


H2O2 


43.48 


19.72 


12 


Can be added to lixiviant 
above ground. Expensive. 


NaC103 


46.97 


21.30 


4.4 


Solubility does not depend 
on pressure. Na can reduce 
permeability. CI can reduce 
ion-exchange resin efficiency. 



1 Assuming $66/kg for U3O8 and neglecting recycling. 



When assessing the significance of 
the chemical costs, it is useful to 
express them in terms of the ppm U3O8 in 
solution that pays for the chemical costs 
of a typical strength lixiviant. Accord- 
ingly, table 2 includes the ppm U3O8 
required to pay for 3 g/1 carbonate or 
bicarbonate, a typical concentration. 
Table 3 lists the ppm U30g required to 
pay for acid molar and normal concentra- 
tions equivalent to 5 g/1 sulfuric acid. 
The costs of equivalent normalities are 
included because they are comparisons of 
the costs of obtaining a selected pH. 
Table 4 lists the ppm U3O8 required to 
pay for 0.3 g/1 oxygen, which is a typi- 
cal concentration, provided by each of 
the oxidizers. A value of $66/kg is 
assumed for U3O8. 



U3O 



The ppm 
were calculated 
the lixiviants, and 
Recycling was not 



in tables 2-4 
recycling of 
so are upper limits, 
included because it 



listed 
assuming no 



depends on site-specific factors. Dis- 
cussions with leaching company personnel 
suggest that 60 to 90 pet of the lixivi- 
ant can be recycled at most sites. The 
ppm values can be compared with the 17 to 
200 ppm U3O8 in the pregnant solutions 
from successful operations. The compari- 
sons can help avoid incorrect conclu- 
sions. For example, one might infer that 
sodium bicarbonate should not be used 
because it costs twice as much as an 
equivalent concentration of dissolved 
carbon dioxide. However, when recycling 
is considered, the cost difference is 
equivalent to only about 1 ppm U3O8 and 
so will have less impact than a very 
small difference in leaching efficiency. 



Tables 2, 3, and 4 also summarize 
the advantages and disadvantages of the 
various alkaline and acid lixiviants and 
the oxidizers, respectively. More 
detailed discussion of each follows. 



carbonate to bicarbonate can be raised by 
adding hydroxide and can be lowered by 
adding carbon dioxide. Leaching compa- 
nies have tried carbonate-bicarbonate 
lixiviants having pH values from 6.5 to 
10. The current trend appears to be to 
use neutral or only moderately alkaline 
lixiviants, avoiding 
except for special 
lower pH values can 
slower dissolution of 



pH values over 9 

circumstances. The 

result in somewhat 

uranium, but they 



minimize the geochemical disturbance to 
the formation, thereby facilitating res- 
toration and usually reducing plugging. 



Compared with acids, al 
iants offer more selectivity 
(that is, dissolve less of 
substances) and can be used 
having a higher carbonate 
calcite) content. In depos 
to both alkaline and acid li 
alkaline lixiviants often 
what less uranium. 



kaline lixiv- 

for uranium 

the unwanted 

in deposits 

(for example, 

its amenable 

xiviants, the 

recover some- 



Ammonium Carbonate-Bicarbonate 

Ammonium carbonate-bicarbonate has 
been the most widely used lixiviant. Its 
advantages include fairly low cost and 
little damage to permeability. However, 
restoration regulations are causing com- 
panies to avoid ammonium lixiviants, 
especially in Wyoming. A Wyoming regula- 
tion, promulgated in 1980, requires that 
the post-restoration concentration of 
ammonium ions be no greater than 0.5 ppm 
if the preleach water was suitable for 
domestic use. Such restoration cannot be 
achieved at a reasonable cost. In fact, 
restoration of ammonium ion levels to 
below 50 ppm is usually expensive and 
requires extended flushing of the ore- 
body. Companies evaluating ammonium 
carbonate-bicarbonate should carefully 
consider the restoration requirements 
that they will have to meet. 



Alkaline (Carbonate Bicarbonate) 
Lixiviants" 

All the currently used alkaline lix- 
iviants contain carbonate-bicarbonate 
that combines with uranium oxide to form 
a soluble anionic complex. The ratio of 



Ammonium carbonate-bicarbonate is 
usually made onsite by combining ammonia, 
water, and carbon dioxide. The chemical 
reactions are as follows: 

2NH3 + H2O + CO2 > (NH4)2 CO3 



NH3 + 



H20 



+ CO. 



> NH4HC03 



The cost of ammonia is about 
$ 0.1 0/1 b not including transportation to 
the site. The cost of carbon dioxide can 
vary considerably, as will be explained 
in the section on carbon dioxide, depend- 
ing on the quantity purchased per year. 
For many operations purchasing 500 to 
2,000 tons per year, it will cost about 
$0. 040/1 b, so this figure was used in 
this report when calculating lixiviant 
costs. Ammonium carbonate then costs 
$11.38/kg-mole, $5.16/lb-mole, $0.12/kg, 
or $0. 054/lb. Ammonium bicarbonate costs 
$7.63/kg-mole, $3.46/1 b-mole, $0.097/kg, 
or $0, 044/1 b. 

The concentration of ammonium 
carbonate-bicarbonate is usually 1 to 
5 g/1 , although high-strength high-pH 
lixiviants containing up to 10 g/1 at 
pH 10 have been employed when attempting 
to prevent calcite dissolution. The 
lower strength lixiviants may dissolve 
uranium more slowly, but they made resto- 
ration less difficult and reduce total 
lixiviant consumption. The significant 
concentration is that of the anion. 
Ammonium carbonate is 62.5 wt-pct carbon- 
ate and ammonium bicarbonate is 77.3 wt- 
pct bicarbonate, so providing 3 g/1 
carbonate or bicarbonate requires 4.8 g/1 
ammonium carbonate or 3.9 g/1 ammonium 
bicarbonate, respectively. These solu- 
tions cost 0.057(^/1 (0.22!^/gal) and 
0.038j[!/l (0.14(z!/gal), respectively. 

When estimating total cost of the 
lixiviants, it would be misleading to 
multiply the number of liters of lixivi- 
ant expected to pass through the forma- 
tion by the cost per liter. The fact 
that most of the lixiviant can be 
recycled by adding a relatively small 
amount of ammonium carbonate-bicarbonate 
lowers total lixiviant consumption con- 
siderably. However, the loading of ion- 
exchange sites can consume significant 
amounts of lixiviant. At many sites, the 
lixiviant cation is exchanged for cal- 
cium, thereby consuming carbonate through 
calcium carbonate precipitation. Thus, 
the exchange capacity can be an important 
factor in projecting costs. Judging 



from discussions with leaching company 
personnel, it would be reasonable to 
assume that about 80 pet of the lixiviant 
can be recycled after the exchange sites 
have been satisfied. 

The reactions with uranium can be 
described as follows: 

UO3 + H2O + 3(NH4)2C03 
^ (NH4)4 UO2 (003)3 + 2NH4OH 

UO3 + 2NH4HCO3 
■> (NH4)2 UO2 (003)2 + H2O 

UO3 + 2NH4HCO3 + (NH4)2C03 

> (NH4)4 U02(C03)3 + H2O 

Sodium Carbonate-Bicarbonate 

The use of sodium carbonate- 
bicarbonate is increasing, as companies 
seek alternatives to ammonium carbonate- 
bicarbonate. Its major drawback has been 
that it tends to swell smectite (montmo- 
rillonite) and plug the formation. 
Sodium has a high hydration energy, so 
water tends to enter clays where sodium 
is absorbed. It is monovalent, so it 
does not hold the clay layers together as 
tightly as calcium and magnesium do. 
Ammonium and potassium ions cause less 
swelling because they fit better between 
the clay layers and have a lower energy 
of hydration. 

Sodium carbonate-bicarbonate can be 
used where the natural sodium content of 
the groundwater is high, or where the 
concentration of swelling clays is very 
low. For example, clay swelling by 
sodium will not be a problem where nearly 
all the clay is a nonswelling type such 
as kaolinite. Companies have had some 
success in keeping clay swelling to tol- 
erable levels, even where those condi- 
tions are not satisfied, by keeping the 
pH near neutral. The reason has not been 
established, but it may be that fewer 
cation exchange sites are occupied by 
sodium as more are occupied by hydrogen. 
This technique is not always successful. 



The advantages of sodium carbonate- 
bicarbonate in formations where it does 
not cause plugging include low cost and 
ease of restoration. A good example of 
its successful use is the OPI-Western 
joint venture pilot operation in Wyoming 
in 1979 (_n ) . Acceptable groundwater 
restoration standards were met after cir- 
culation of only six pore volumes of 
surface-treated water through the mined 
aquifer. The lixiviant just before the 
start of restoration contained 1.0 g/1 
sodium and 0.9 g/1 bicarbonate. 

Sodium carbonate lixiviant is usu- 
ally made by dissolving soda ash (anhy- 
drous sodium carbonate) or by bubbling 
carbon dioxide through caustic soda 
(sodium hydroxide solution). In Wyoming, 
soda ash is cheaper. In Texas, using 
caustic soda may be cheaper because of 
freight charges on soda ash shipped from 
Wyoming. Sodium carbonate from soda ash 
costs about $10.52/kg-mole, $4.77/lb- 
mole, S0.099/kg, or S0.045/lb, not 
including transportation to the site. It 
costs considerably more when made from 
caustic soda and carbon dioxide, 
S47.98/kg-mole, $21.76/lb-mole, S0.45/kg, 
or S0.21/lb. 



Sodium bic 
ash and carbon 
mole, $3.27/1 
S0.039/lb. Whe 
ide and carb 
$25.93/kg-mole, 
or $0.14/lb. 
or powder can b 
mole, 310.08/ 
$0.12/lb. 



arbonate made from soda 
dioxide costs $7.19/kg- 
b-mole, $0.086/kg, or 
n made from sodium hydrox- 
on dioxide, it costs 
$11.76/lb-mole, $0.31/kg, 
Sodium bicarbonate flakes 
e purchased for $22.22/kg- 
Ib-mole, S0.26/kg, or 



Sodium carbonate-bicarbonate can 
also be made from purified trona, mar- 
keted as sodium sesquicarbonate (Na2C03 
• NaHC02 • 2H2) (9_). Trona is mined pri- 
marily in Wyoming and is processed to 
make soda ash. Even though less process- 
ing is required to make sodium sesqui- 
carbonate, its cost per ton in 1980 was 
similar to that of soda ash because of 
the smaller market. Using the prices in 
table 1, sodium carbonate made by combin- 
ing sodium sesquicarbonate and caustic 
soda costs $17.60/kg-mole, $7.98/lb-mole. 



$0.17/kg, or 

ate made by 
through sodi 
costs $8.6 
$0.10/kg, or 
market devel 
may provide 
carbonate-bic 
readily than 
mixing costs. 
nient for 
with a lixiv 
concentration 
ate because n 
to be added. 



$0.075/lb. 
bubbl ing 
um sesquica 
7/kg-mole, 
$0.047/lb. 
ops, sodium 
a cheaper s 
arbonate. I 
soda ash, 
It would 
perators wh 
iant contain 
s of carbona 
carbon di 



Sodium bicarbon- 

carbon dioxide 

rbonate solution 

$3.93/lb-mole, 

If a sufficient 

sesquicarbonate 

ource of sodium 

t dissolves more 

thereby reducing 

be yery conve- 

wish to leach 

ing equal molar 

te and bicarbon- 

oxide would have 



The concentration of sodium 
carbonate-bicarbonate in a lixiviant is 
usually 1-5 g/1, just as for ammonium 
carbonate-bicarbonate. A pH up to 9 or 
10 has been used, but the trend now 
appears to be to keep the pH near neutral 
to minimize clay swelling. Sodium car- 
bonate is 56.6 wt-pct carbonate and 
sodium bicarbonate is 72.6 wt-pct bicar- 
bonate, so providing 3 g/1 carbonate or 
bicarbonate requires 5.3 g/1 sodium car- 
bonate or 4.1 g/1 sodium bicarbonate, 
respectively. If made using soda ash, 
these solutions cost 0.053?!/l (0.20?;/gal) 
and 0.035(2f/l (0.13«f/gal ) , respectively. 
If made using caustic soda, they cost 
0.24?:/l (0.91izf/gal) and 0.13?f/l 
(0.48?!/gal ) , respectively. If made by 
combining sodium sesquicarbonate and 
caustic soda or carbon dioxide they cost 
0.090^/1 (0.34?:/gal) and O.Q^U/I 
(0.162f/gal ) , respectively. Thus, sodium 
carbonate-bicarbonate made using soda ash 
may be slightly cheaper than ammonium 
carbonate-bicarbonate. The difference in 
total chemical costs will be enhanced by 
the fact that less sodium will be 
absorbed on clays, so the consumption of 
a sodium lixiviant will be smaller than 
that of an ammonium lixiviant. 

The chemical reactions of oxidized 
uranium with sodium carbonate-bicarbonate 
are similar to those with ammonium 
carbonate-bicarbonate. For example, the 
chemical equation corresponding to the 
equation under ammonium carbonate- 
bicarbonate is as follows: 



UOj + Na2C03 + 2NaHC03 
^ Na4 [U02(C03)3]+ H2O. 

Potassium Carbonate-Bicarbonate 

Potassium carbonate-bicarbonate is 
very attractive technically and environ- 
mentally. It does not swell clays as 
sodium does. Meeting restoration 
requirements should be less difficult 
than with ammonium lixiviants. Potassium 
is normally found in the preleach ground- 
water, so its presence after leaching 
does not change the potential uses of the 
water. However, its high cost has pre- 
vented its use. 

Potassium carbonate can be made by 
bubbling carbon dioxide through caustic 
potash (potassium hydroxide solution). 
Made in this way, it costs $63.26 kg- 
mole, $28.69/lb-mole, $0.46/kg, or 
$0.21/1 b. A more convenient but somewhat 
more expensive source is a liquid con- 
taining 47 pet potassium carbonate. It 
then costs $76.18/kg-mole, $34.55/lb- 
mole, $0.55/kg, or $0.25/lb. (All costs 
are on a 100 pet basis.) Potassium 
bicarbonate made by bubbling additional 
carbon dioxide through caustic potash 
costs $33.56/kg-mole, $15.22/lb-mole, 
$0.34/kg, or $0.15/lb. Potassium bicar- 
bonate granules can be purchased for 
$30.90/kg-mole, $14.00/lb-mole, $0.31/kg, 
or $0.14/lb. If an operator wished to 
avoid handling carbon dioxide, the potas- 
sium bicarbonate granules could be com- 
bined with the 47 pet potassium carbonate 
in the proportion to give any desired 
pH. 

A new technique developed by the 
University of Texas at Austin under 
Bureau of Mines research contract 
H0282016 promises to greatly reduce the 
cost of using potassium carbonate- 
bicarbonate and also reduce permeability 
loss from calcium carbonate precipita- 
tion. The technique involves flushing 
the formation with potassium chloride 
before injecting potassium carbonate- 
bicarbonate. This satisfies the cation 
exchange sites with potassium from rela- 
tively cheap potassium chloride, which 
costs about $4.92/kg-mole, $2.24/lb-mole. 



$0.066/kg, or $0.030/lb. In a laboratory 
column leaching experiment, the consump- 
tion of potassium carbonate following 
chloride preflush was only 17 pet of that 
without the chloride preflush. The pre- 
flush reduces permeability loss because 
soluble calcium chloride is formed 
instead of insoluble calcium carbonate. 
The calcium can be removed aboveground 
before injecting potassium carbonate. 
The problem of what to do with calcium 
aboveground is not negligible, especially 
since the calcium solution will contain 
some radium. However, the problem of a 
plugged formation can have a much more 
severe impact on a leaching operation 
than the calcium disposal. The chloride 
preflush is more fully described in 
another publication (35). 

The desired pH range and concentra- 
tion of carbonate-bicarbonate for potas- 
sium carbonate-bicarbonate should be sim- 
ilar to those for ammonium or sodium 
carbonate-bicarbonate. An operator may 
have the option of using a higher pH and 
concentration because the problems of 
high-pH, high-strength ammonium lixiviant 
creating restoration difficulties, and 
high-pH, high-strength sodium lixiviant 
creating clay swelling difficulties are 
both avoided. Potassium carbonate is 
43.4 wt-pct carbonate and potassium 
bicarbonate is 60.9 wt-pct bicarbonate, 
so providing 3 g/1 carbonate or bicar- 
bonate requires 6.9 g/1 potassium carbon- 
ate or 4.9 g/1 potassium bicarbonate, 
respectively. If made using caustic pot- 
ash, these solutions cost 0.32j^/l 
(1.20iz!/gal) and 0.16izl/l (0.62j[l/gal ) , 
respectively. The carbonate solution 
made from 47 pet potassium carbonate liq- 
uid costs 0.38j[!/l (1.43(2f/gal). The 
bicarbonate solution made from granules 
costs 0.15jz!/l (0.57jz:/gal). 

The chemical reactions of oxidized 
uranium with potassium carbonate- 
bicarbonate are similar to those with 
ammonium carbonate-bicarbonate. 

Carbon Dioxide 

Lixiviants made by bubbling carbon 
dioxide through groundwaters from the 



10 



formation to be leached offers technical, 
economic, and environmental advantages. 
It is the cheapest lixiviant on a molar 
basis. It creates the least geochemical 
disturbance in the formation and should 
not cause clay swelling. It should allow 
easier restoration than any other lixivi- 
ant. Its use is increasing; it is being 
used in at least two commercial 
operations. 

This lixiviant does not leach effec- 
tively in all formations, however. Dis- 
cussions with leaching companies indicate 
that it is successful in deposits having 
clean sands, but does not leach effec- 
tively if much organic carbon is present. 
Unpublished laboratory data indicated 
that it leached the Wyoming ore being 
tested somewhat slower than ammonium 
carbonate-bicarbonate. One leaching com- 
pany employee felt that it can leach 
faster than ammonium carbonate- 
bicarbonate in some deposits and usually 
leaches almost as fast in amenable 
deposits. 

The lixiviant produced by bubbling 
carbon dioxide through the groundwater 
should not be thought of as carbonic 
acid. Dissolved carbon dioxide is most 
effective where there is enough carbonate 
in the groundwater to serve as a buffer 
and form bicarbonate. Then the reactions 
are primarily of the form 



CO^ 



Na2C03 



H2O 



2NaHC0: 



and 



CO2 + K2CO3 + H2O ^ 2KHCO3 

rather than 

CO2 + H2O ->■ H2CO3. 

The pH of this lixiviant is usually 6.5 
to 7.5. 

The cost of carbon dioxide varies a 
great deal, depending primarily on the 
quantity purchased per year, and also on 
the distance from the seller and on local 
supply-demand factors. The prices listed 
by one seller in Texas are as follows: 



Tons purchased per year Cost, $/ton 

Over 2,000 61 

1,001-2,000 77 

501-1,000 96 

301-500 125 

101-300 130 

51-100 195 

1-50 230 

These prices include delivery within 
50 miles. Transportation charges add 
$10/ton for each additional 50 miles. 
The prices quoted by all sellers simi- 
larly depended on quantity purchased per 
year. However, for a given quantity, the 
price varied more with location than it 
did for most other chemicals. 

To determine a typical price to use 
when calculating costs of lixiviants it 
is necessary to choose a typical consump- 
tion per year. For most operations, the 
ratio of carbon dioxide consumed to U3O8 
produced will be between 5 and 10 on a 
weight basis. Thus, a large commercial 
operation producing 500,000 to 1 mil- 
lion lb U3O8 per year might consume 
2,000 tons of carbon dioxide per year. 
Most commercial operations would consume 
over 500 tons per year. Discussions with 
carbon dioxide sellers and in situ leach- 
ing companies indicated that $80/ton 
would be a typical price for consumption 
fo 1,000 to 2,000 tons per year at most 
locations in Wyoming and Texas. Using 
$80/ton, the cost of carbon dioxide is 
$3.88/kg-mole, $1.76/lb-mole, $0.088/kg, 
or $0. 040/1 b. 

Ideally, the molar concentration of 
bicarbonate in this lixiviant should be 
similar to that in ammonium or sodium 
bicarbonate. In practice, dissolving 
that much carbon dioxide without exces- 
sive lowering of the pH can seldom be 
done, so the concentration is usually 
lower. However, useful cost comparisons 
can be made assuming equivalent concen- 
trations. Each mole of carbon dioxide 
can provide a mole of bicarbonate, assum- 
ing that all the carbon dioxide dissolves. 
Thus, the cost of solution made from car- 
bon dioxide providing 3 g/1 bicarbonate 
(2.2 g/1 carbon dioxide) is only 0.019^/1 



11 



(0.073(^/gal), about half that of the sec- 
ond cheapest lixiviant. The difference 
in total chemical costs will be even 
greater because this lixiviant should 
have smaller consumption by cation 
exchange than any other. 

To obtain the potential cost savings 
from using carbon dioxide, it is vital 
that an efficient sparging system produc- 
ing yery small bubbles be employed for 
dissolving the carbon dioxide in the lix- 
iviant. A system that is not carefully 
designed and operated can easily waste 
half the carbon dioxide by allowing it to 
pass through the solution as large bub- 
bles that escape before dissolving. This 
is especially true as the pH is brought 
down to near neutral. An additional cost 
factor is that storing carbon dioxide in 
bulk as a liquefied gas requires a spe- 
cial tank and equipment because it must 
be cooled during storage. 

The chemical reactions with oxidized 
uranium are similar to those with ammo- 
nium bicarbonate. 

Acid Lixiviants 

Acid lixiviants are used much less 
than alkaline lixiviants. Many (perhaps 
most) deposits are not amenable to acid 
leaching because they contain too high a 
concentration of acid consumers, usually 
calcium carbonate. In amenable deposits, 
acid lixiviants generally recover some- 
what more uranium than alkaline lixivi- 
ants do. They also dissolve more of most 
of the other elements, especially metals, 
so the total dissolved solids are higher. 
Sulfuric acid has been used successfully. 
Hydrochloric and nitric acids have been 
proposed for special applications, but 
are not being used. 

Sulfuric Acid 

Sulfuric acid leaches very effec- 
tively in amenable deposits. Its techni- 
cal advantages and disadvantages both 
stem from the fact that it dissolves more 
gangue than alkaline lixiviants do. Sul- 
furic acid often can recover more uranium 
than alkaline lixiviants, based on pub- 



lished (10) and unpublished laboratory 
data. TFe higher recovery may be due to 
more of the uranium being exposed to the 
lixiviant, as more of the gangue is 
dissolved. 

The acid dissolves higher concentra- 
tions of some pollutants such as thorium 
and toxic metals, but less of radium and 
selenium. Restoration after sulfuric 
acid leaching at the Rocky Mountain 
Energy Co.'s Nine Mile operation (7^) in 
Wyoming required extensive flushing but 
was successful. It appears that restor- 
tion after sulfuric acid is more 
difficult than after sodium carbonate- 
bicarbonate, but easier than after ammo- 
nium carbonate-bicarbonate. 

There is disagreement as to whether 
harmful channeling occurs with acid 
leaching. Theoretically, acid could dis- 
solve the gangue to form pathways for the 
lixiviant that bypass most of the ore. 

Concentrated sulfuric acid (96- 
98 pet) costs about $8.65/kg-mole, 
$3.92/lb-mole, $0.088/kg, or $0.040/lb. 
A pH of 1.8 to 2.0 is preferred. Unlike 
carbonate-bicarbonate lixiviants, the pH 
cannot be controlled independently of the 
concentration. Theoretically, obtaining 
a pH of 2 in distilled water would 
require only 0.5 g/1 H2SO4. At the Rocky 
Mountain Energy Co.'s operation, about 
5 g/1 were required to obtain a pH of 2. 
A 5 g/1 solution costs about 0.044(^/1 
(0.17(Zl/gal). 

When considering the amenability of 
a deposit to acid leaching, both the 
cost of the consumed acid and the pos- 
sible blockage from calcium sulfate and 
from carbon dioxide gas should be consid- 
ered. Calculation of the cost of the 
acid consumed by calcite is straightfor- 
ward if the average calcite concentration 
is known. For example, assume that the 
deposit is 1.0 pet calcite. Because 98 g 
H2SO4 is required to consume 100 g CaC03, 
the calcite in each kilogram of ore will 
consume (0.01) (98/100) = 0.0098 kg 
H2SO4, costing 0.086^. If the ore con- 
tains 0.05 pet recoverable uranium worth 
$66/kg, then 1 kg of ore yields uranium 



12 



worth 3.3^. Thus, if using acid 
increases uranium recovery by 3 pet, the 
value of the extra uranium recovered 
exceeds the cost of the acid consumed by 
the calcite in this example. This exam- 
ple suggests that a calcite concentration 
of 1 or 2 pet does not by itself lead to 
prohibitively high costs for acid con- 
sumption. Some acid will be consumed by 
other components in the ore, of course. 

Blockage by calcium sulfate or car- 
bon dioxide can occur if their solubility 
is exceeded. The problem can be mini- 
mized by increasing the acid strength 
slowly, so that the products are formed 
slowly and lixiviant can transport them 
to a producTtion well without the solubil- 
ity being exceeded. Operators familiar 
with acid leaching have stated that cal- 
cium sulfate is a more serious problem 
than carbon dioxide. Carbon dioxide 
forms only when the acid dissolves cal- 
cite or other carbonates. It has a high 
solubility under typical formation pres- 
sures. Calcium sulfate can form when 
calcium is removed from minerals by ion- 
exchange, in addition to forming when 
calcite is dissolved. 

Although ionic strength affects sol- 
ubilities, the published solubilities in 
water provide a guide as to the approxi- 
mate conditions under which solubilities 
in lixiviants might be exceeded. The 
solubility of calcium sulfate in water at 
18° C is 2.0 g/1 (28, p. 249). To calcu- 
late the concentration of calcite in the 
ore that could produce a saturated solu- 
tion of calcium sulfate by reacting with 
a stoichiometric amount of sulfuric acid, 
assume that the porosity of the ore is 
33 pet and that the density of the rock 
particles is 2.6 g/cm^. Then, to produce 
saturation, 100 cm^ of ore must contain 
sufficient calcite to produce 
(0.33 1) (2.0 g/1) = 0.66 g of calcium 
sulfate. The corresponding amount of 
calcite is (0.66) (100/136) = 0.49 g, 
which gives a calcite concentration of 
(0.49)/(667)(2.6) or 0.03 pet. A greater 
than stoichiometric concentration of sul- 
furic acid would allow the solubility 
product of calcium sulfate to be exceeded 
with an even lower concentration of 
calcite. These calculations show that in 



nearly all formations, rapid introduc- 
tion of strong acid could produce cal- 
cium sulfate greatly in excess of its 
solubility. 

Similar calculations show that more 
calcite is required to exceed the solu- 
bility of carbon dioxide, but blockage 
could be a problem if strong acid is rap- 
idly injected, especially into shallow 
deposits. Consideration of published 
data (28, pp. 365-368) indicates that in 
the range of 170 to 680 ft of water 
absolute pressure (5 to 20 atmospheres) 
and 15** to 25° C, the solubility (S) in 
grams per liter of carbon dioxide in 
water can be approximated by 

S = 9.6 + 0.044(P-170) - 0.0012P(T-15) , 

where P is the absolute pressure in feet 
of water and T is the Celsius tempera- 
ture. (Extrapolating outside the stated 
range of pressure and temperature will 
produce large errors because the experi- 
mental solubility is a nonlinear function 
of pressure and temperature.) Within the 
stated range, the lowest solubility is at 
170 ft of water and 25° C, where 
S = 7.6 g/1 and the highest is at 680 ft 
of water and 15°, where S = 32 g/1. Cal- 
culations similar to those for calcium 
sulfate show that when S = 7.6 g/1, the 
solubility of the carbon dioxide from 
calcite could be exceeded if the ore is 
greater than 0.3 pet calcite. When 
S = 32 g/1, the solubility could be 
exceeded if the ore is greater than 
1.4 pet calcite. 

Sulfuric acid reacts with oxidized 
uranium to form soluble anionic complexes 
of sulfate and uranium dioxide, according 
to the following equations. 

UO3 + H2SO4 ->- UO2SO4 + H2O 

UO2SO4 + H2SO4 ^ H2 (UO2 (S04)2) 

H2 (UO2 (S04)2) + H2SO4 > H4 (UO2 ($04)3) 

Nitric and Hydrochloric Acids 

Nitric and hydrochloric acids are 
not being used for in situ leaching but 
have been proposed for leaching ore that 



13 



cannot be leached by other lixiviants. 
Nitric acid offers the advantage of being 
an oxidizer, and both acids are said to 
leach more effectively than sulfuric acid 
in ores containing much shale, clay, or 
carbonaceous material. A disadvantage is 
that they do not form anionic complexes 
with uranium. Thus, while uranium can be 
extracted from sulfuric acid with highly 
selective anionic exchange resins, ura- 
nium must be extracted from nitric or 
hydrochloric acids with less selective 
cationic exchange resins (19^, p. 60). 

They have environmental disadvan- 
tages. They share the nonselectivity of 
sulfuric acid and in addition will read- 
ily dissolve radium. Radium nitrate and 
chloride are much more soluble than 
radium sulfate or bicarbonate (5^). In 
addition, nitric acid will almost cer- 
tainly be subject to more restrictive 
restoration requirements than sulfuric or 
hydrochloric acid. The EPA drinking 
water standards for sulfate and chloride 
are both 250 ppm, but the standard for 
nitrate nitrogen is only 10 ppm. (The 
limit is expressed in terms of nitrogen 
attributable to nitrate because standard 
tests measure nitrogen; 10 ppm nitrogen 
from nitrate is equivalent to 44 ppm 
nitrate.) 

Comparing the cost of nitric and 
hydrochloric acids with that of sulfuric 
acid is complicated by the difference in 
reaction mechanisms. Perhaps the most 
useful comparison is on the basis of nor- 
mality, which is equivalent to comparing 
the costs for producing a given pH. 
Using that basis, nitric and hydrochloric 
acids cost about twice as much as sul- 
furic acid. Concentrated nitric acid 
(70 pet) costs about $8.33/kg-mole, 
$3.78/lb-mole, $0.13/kg, or $0.060/lb on 
a 100 pet HNO3 basis. Concentrated 
hydrochloric acid (37 pet) costs about 
$0. 035/lb, and so costs $7.61 kg-mole, 
$3.45/lb-mole, $0.21/kg, or $0.095/lb on 
a 100-pct basis. Solutions of 0.102 ^ 
(the same normality as 5 g/1 sulfuric 
acid) nitric and hydrochloric acids cost 
0.085(z!/l (0.32(Zl/gal) and 0.078j[!/l 
{0.29^/gal), respectively. The total 
chemical cost when using nitric acid may 



be reduced by savings in oxidizer. 

Oxidizers 

The oxidizers currently used for in 
situ uranium leaching are oxygen and 
hydrogen peroxide. Chlorates have been 
used and are discussed for comparison. 

Oxygen 

Oxygen is used at most commercial - 
scale operations because of its low cost. 
It is usually shipped and stored as a 
liquid in bulk, although pilot operations 
have used compressed oxygen in bottles. 
In bulk, it costs about $3.40/kg-mole, 
$1.55/lb-mole, $0.11/kg, or $0.048/lb in 
Texas and $5.10/kg-mole, $2.33/lb-mole, 
$0.16/kg, and $0. 073/1 b in Wyoming. Typ- 
ical concentrations in lixiviant are 0.1 
to 0.3 g/1. In Texas 0.3 g/1 costs 
0.0033)^/1 (0.012!^/gal), and in Wyoming 
costs 0.0048(z!/l (0.018jZl/gal ) . Although 
this is a lower cost per liter than the 
other components in the lixiviant, the 
total cost of the oxidizer can be greater 
because so little of it is recycled. 

The disadvantages of oxygen are 
caused by its limited solubility. Litz 
( 16 ) reported that the solubility (S) in 
grams per liter of oxygen in lixiviants 
could be approximated by 

°*°^^ ^ (1.107 - 0.07 log P), 



S = 



33.5 + T 



where P is the absolute 
of water and T is the 
ture. The formula in 
solubility under atmo 
(33.9 ft of water) at 
0.04 g/1, and that 300 
required to produce a 
0.3 g/1 . Thus, oxygen 
under pressure to obtain 
bility. It is usually 
at the leaching depth, 
a separate oxygen line 
each injection well. 



pressure in feet 

Celsius tempera- 

dicates that the 

spheric pressure 

20° C is only 

ft of water is 

solubility of 

must be injected 

sufficient solu- 

injected downhole 

thereby requiring 

and sparger for 



An additional disadvantage is the 
risk of bubbles forming and causing gas 
phase blockage. Bubbles may form even 



14 



when the solubility is not exceeded 
because thorough mixing of the oxygen and 
lixiviant is not instantaneous. Small 
bubbles can form in and close to an 
injection well. These bubbles do not 
pass between sand grains as lixiviant 
flows into the formation, so the bubbles 
can build up and cause some blockage even 
when the theoretical solubility is not 
exceeded. An efficient sparger that pro- 
duces very small bubbles helps to mini- 
mize such blockage. 

Hydrogen Peroxide 

Hydrogen peroxide has been used for 
the majority of pilot scale operations. 
It is convenient because it can be added 
to the lixiviant at one central point 
under atmospheric pressure instead of 
requiring separate lines to all injection 
wells. A higher concentration of hydro- 
gen peroxide than oxygen can be injected 
without causing gas phase blockage. This 
advantage is especially important for 
shallow deposits. 

Studies have indicated that hydrogen 
peroxide oxidizes more effectively than 
oxygen (2j^). However, there is disagree- 
ment as to the extent to which this 
advantage assists in the field. Most 
investigators believe hydrogen peroxide 
decomposes to oxygen so rapidly that it 
is not superior to an equivalent amount 
of oxygen. A differing opinion is that 
even though most of the hydrogen peroxide 
decomposes, much of the ore is contacted 
by enough hydrogen peroxide to affect the 
leaching rate. 

The decomposition is slower in acid 
than in alkaline lixiviants. Hydrogen 
peroxide is much more stable in acid than 
in alkaline solutions of the same purity 
(8^). Researchers familiar with hydrogen 
peroxide stated that the effect of alka- 
linity would be greater than the catalyz- 
ing effect of the higher concentration of 
dissolved metals in acid lixiviants. 



Hydrogen peroxide for lixiviants is 
purchased and transported in solutions of 
either 50 or 70 pet H2O2. Because of 
requirements of the National Fire Protec- 



tion Association, it is usually stored 
onsite at less than 51 pet (12) e Ship- 
ping it as a 70 pet solution can save 
some freight charges, but then facilities 
for diluting the solution with high qual- 
ity water must be provided. 

In August 1980, 50 pet hydrogen per- 
oxide cost 25.75^/1 b and 70 pet cost 
36.00j^/lb, f.o.b., 3,500 gal minimum 
(12). The freight cost to most locations 
would be 3(^/1 b to 5^/1 b on an "as-is" 
basis. This report uses 58^/1 b on a 100- 
pct basis as a representative price. 
Then the • cost of hydrogen peroxide is 
$43.48/kg-mole, $19.72/lb-mole, $1.28/kg, 
or $0.58/1 b on a 100-pct basis. Each 
mole of hydrogen peroxide provides half a 
mole of oxygen, and each kilogram of 
hydrogen peroxide provides (16/34) kg of 
oxygen, so the effective cost of oxygen 
purchased as hydrogen peroxide is 
$86.96/kg-mole, $39.44/lb-mole, $2.72/kg, 
or $1.23/lb. Typical concentrations of 
hydrogen peroxide in lixiviants are 0.3 
to 1.0 g/1 , which would provide the same 
oxygen as 0.14 to 0.47 g/1 O2. The cost 
of 0.3 g/1 oxygen supplied by hydrogen 
peroxide is 0.082iz!/l (0.31jz!/gal ). 

The high cost of hydrogen peroxide, 
coupled with the fact that little can be 
recycled, results in it being the major 
chemical cost at most operations using 
it. To illustrate, if the hydrogen per- 
oxide costs 0.082(^/1 of lixiviant and 
uranium is worth $66/kg, then 12 ppm U3O8 
in the pregnant lixiviant is required to 
pay for the oxidizer. 

Chlorates 

Sodium chlorate has been used as an 
oxidizer (32), and potassium chlorate 
could be used. Their cost is between 
those of hydrogen peroxide and oxygen.. 
Sodium chlorate costs about $46,97/kg- 
mole, $21.30/lb-mole, $0.44/kg, or 
$0.20/lb. Potassium chlorate costs only 
4 pet more on a molar basis. Considering 
the fact that each molecule of sodium 
chlorate provides 3 atoms of oxygen shows 
that the cost of oxygen from this source 
is $31.31/kg-mole, $14.20/lb-mole, 

0.98/kg, or $0.44/lb. Providing 0.3 g/1 



15 



oxygen would cost 0.029jzl/l (O.lljzl/gal ). 
The total cost of chlorate relative to 
the other oxidizers may be reduced some- 
what by the ability to recycle unconsumed 
chlorate, whereas unconsumed oxygen goes 
out of solution in the surface plant and 
is lost. 

Chlorates are effective oxidizers in 
acid solutions that contain iron (19, 
p. 64). Apparently, the chlorates oxi- 
dize iron and the iron oxidizes the ura- 
nium. However, neither sodium nor 
potassium chlorate oxidizes efficiently 
in carbonate-bicarbonate lixiviants (19, 
p. 104). One researcher, discussing 
unpublished experiments, stated that 
chlorates in carbonate-bicarbonate lixiv- 
iants oxidized uranium to some extent but 
were less effective than oxygen. 

Other disadvantages result from the 
buildup of chloride as the lixiviant is 
recycled repeatedly. The chloride 
increases the corrosiveness of the lixiv- 
iant and decreases the uranium loading 
capacity of ion-exchange resin. Some 
companies believe that the chloride would 
make restoration more difficult. The 
sodium in sodium chlorate could reduce 
permeability. These problems do not com- 
pletely rule out the use of chlorates. 
They could be advantageous in very shal- 
low deposits, for example, especially 
with acid lixiviants. 

METHODS OF TESTING LIXIVIANTS 

The costs, advantages, and disadvan- 
tages previously presented provide only a 
general guide for lixiviant selection. 
To determine the suitability for a spe- 
cific deposit, thorough laboratory and 
field testing is necessary. 

Laboratory Tests 

Both batch leach tests (sometimes 
called agitation leach tests) and column 
leach tests are used in selecting the 
lixiviant. Batch leach tests consist of 
placing the ore and lixiviant in a con- 
tainer, often a sealed flask, and gently 
agitating them. Column leach tests 
consist of passing lixiviant through a 
column packed with ore. 



Batch Leaching 

Although batch leaching tests do not 
simulate downhole conditions, they do 
provide useful relative data. They show 
the relative rate and amount of uranium 
extraction with tested lixiviants and can 
give an indication of lixiviant and oxi- 
dant consumption. 

The ore sample must be large enough 
to allow solution samples to be withdrawn 
without unduly disturbing the geochemis- 
try and to provide sufficient dissolved 
uranium for analysis. Typical ore sample 
sizes are 100 to 500 g, with a liquid- 
solid volume ratio of three. 

The variables are usually type of 
lixiviant (for example, sodium or ammo- 
nium bicarbonate), pH (for example, 7, 8, 
and 9), carbonate or bicarbonate concen- 
tration (for example 1, 3, and 5 g/1), 
and oxidizer concentration (for example, 
equivalent to 0.1, 0.3, and 0.5 g/1). 
When performing tests to select a lixivi- 
ant for a particular site, the solutions 
should be made from groundwater from the 
formation to be leached, not from dis- 
tilled water. (Simulated groundwater 
made by adding measured quantities of 
chemicals to distilled water may be pre- 
ferred for general studies of reaction 
mechanisms.) The ore should be blended 
to ensure homogeneity, but not ground. 

The measured observables are usually 
pH and concentrations of uranium, carbon- 
ate and/or bicarbonate, oxidizer, and 
site-specific elements such as vanadium 
that may be present in sufficient concen- 
tration to interfere. Some experimenters 
measure Eh, others feel that Eh measure- 
ments in this type of experiment are not 
meaningful. Typical sampling times are 
0, 1, 4, 8, 12, and 24 hours after the 
start of the test and once per day until 
equilibrium is reached, usually in 3 days 
or less. The uranium in the ore should 
be measured before and after leaching and 
a material balance made to check the 
validity of the measurements. There is 
disagreement among experimenters as to 
whether fresh lixiviant should be added 
to replace lixiviant withdrawn for sam- 
pling. Either method can give good 



16 



results. Replacing the lixiviant allows 
a smaller sample to be used but compli- 
cates the material balance. 

Measuring the consumption of oxygen 
or hydrogen peroxide requires sealed 
experimental equipment and a method for 
withdrawing samples without losing oxy- 
gen. Some experimenters favor a chlorate 
oxidizer to avoid the need for pressur- 
ization. The consumption of chlorate can 
be determined by measuring the increase 
in chloride. Because the oxidizing 
effectiveness of chlorate may be differ- 
ent from that of oxygen, the consumption 
of chlorate will not necessarily be equal 
to that of oxygen. Nevertheless, oxida- 
tion tests with chlorates can be a useful 
and economical method of comparing the 
amounts of oxidizer consumed by different 
samples. 

Obtaining meaningful results from 
oxidizer consumption tests requires care. 
The results of a careless experiment can 
range from much less to much more than 
the correct value. If the ore is inad- 
vertently oxidized before being leached, 
then the results will be too low. Avoid- 
ing preoxidation of the ore requires that 
the ore be protected as soon as it is 
cored until it is placed in the leaching 
flask. The ore should be sealed in air- 
tight cylinders in nitrogen or frbzen in 
dry ice. Sealing it in plastic is not 
adequate. Wrapping it in foil and seal- 
ing it in wax may be adequate if done 
carefully and the storage time is short. 
A plexiglass cylinder that can be flushed 
with nitrogen after inserting the core 
and then sealed airtight for transporting 
has worked well for the Bureau of Mines. 
If the ore is not preoxidized, then the 
results will provide an upper limit to 
the oxidizer consumption. The results 
are an upper limit because the oxidizing 
reactions will be more complete in an 
agitated flask with a liquid-solid volume 
ratio of three than in actual in situ 
leaching conditions. Some commercial 
service laboratories feel they have 
reliable methods of predicting what the 
field consumption of oxygen will be. 
Those methods are proprietary. 



X-ray diffraction measurements of 
clay swelling can indicate which lixivi- 
ants cause the most swelling and hence 
allow qualitative predictions of permea- 
bility loss. Usually, however, permea- 
bility loss is studied with column leach- 
ing experiments. 

Column Leaching 

Column leaching tests simulate field 
conditions more closely than batch tests, 
but caution must still be used when 
extrapolating from laboratory to field. 
The contact between ore and lixiviant is 
more complete than in actual in situ 
leaching. Therefore, the measured con- 
sumption of lixiviant and oxidizer and 
the extraction of uranium should be 
viewed as upper limits of what might be 
expected in the field. 

Column leach tests can indicate per- 
meability losses, but to obtain meaning- 
ful results, water from the formation 
should be used and the ore should be 
disaggregated and blended. Making lixiv- 
iants by adding only the primary compo- 
nents (for example, ammonium carbonate) 
to distilled water will cause mislead- 
ingly large permeability losses during 
column leaching. Attempts to use intact 
cores in hopes of better simulating down- 
hole conditions have not been satisfac- 
tory. Meaningful comparisons of lixivi- 
ants require similar cores, but cores 
vary considerably in permeability and 
uranium content. Further, the lixiviant 
flows through the core in the ore's ver- 
tical direction but flows horizontally, in 
situ. Therefore, a thin clay lens can 
drastically reduce the measured permea- 
bility of a core but have very little 
effect on the in situ horizontal permea- 
bility. The ore should be disaggregated 
gently to avoid breaking grains and thor- 
oughly blended. Most sandstone uranium 
ore can be disaggregated easily. 

To obtain meaningful data on the 
effect and the consumption of oxidizer, 
the ore must be protected from preoxida- 
tion, just as for batch tests. The type 
of oxidizer and the formation of bubbles 



17 



can affect permeability, so ideally the 
column would be pressurized to simulate 
downhole conditions. Using sodium chlo- 
rate in these tests could reduce 
permeability. 

There is disagreement as to the best 
column configuration. For nonpressurized 
systems, the authors prefer a vertical 
column with lixiviant entering at the 
bottom. Admitting the lixiviant at the 
bottom maximizes the speed at which any 
air or other gases in the core at the 
beginning of leaching will be swept out. 
Horizontal tubes present the risk that 
ore will slump slightly, thereby creating 
a high permeability channel between the 
top of the ore and the inside of the col- 
umn. Some experimenters favor horizontal 
columns. They are more convenient, 
especially if long columns are used. 

There is also disagreement as to the 
choice of diameter and length. Theoreti- 
cally, the most reliable results will be 
obtained with the widest and longest col- 
umns. Greater length better simulates 
chromatographic effects (differing migra- 
tion speeds of different ions). Greater 
diameter minimizes edge effects; in par- 
ticular, the tendency for the flow resis- 
tance to be smaller between the ore and 
the column than in the ore. In practice, 
the amount of core is usually quite lim- 
ited and the number of tests to be run 
quite large, so rather small columns must 
be used. A column 5 cm in diameter and 
120 cm long appears to be a good compro- 
mise unless chromatographic effects are 
of special interest. Then a 2.5 cm diam- 
eter sectional tube with a total length 
of 300 to 500 cm would provide more 
information. In one unpublished experi- 
ment, 1.6-cm-ID horizontal columns were 
used with good results by employing care- 
ful packing techniques. There is some 
disagreement as to the need for simulat- 
ing chromatographic effects when select- 
ing a lixiviant. 

For testing oxidizer consumption and 
response to oxygen under downhole condi- 
tions, a smaller pressure cell can be 
used. Hassler pressurized cells 1.75 cm 
in diameter and 10 cm long have given 



good results at the University of Texas. 
With such small cells, it is important 
that the sides of the cell press tightly 
against the ore to avoid large edge 
effects. 

Typical variables, observables, and 
sampling times are similar to those for 
batch tests. In addition, permeability 
can be measured either by maintaining a 
constant pressure and measuring the flow 
rate or maintaining a constant flow with 
a positive displacement pump and measur- 
ing the pressure difference across the 
column. The second method is generally 
preferred. The speed of the lixiviant 
should be similar to that in the field, 
often about 3 m/day. Excess flow rate 
can lead to channeling, especially with 
horizontal columns. The pump should not 
pulse the flow, as excessive pulses can 
increase the movement of fines and reduce 
permeability. 

Pilot Field Tests 

A pilot-scale field test is essen- 
tial before starting commercial opera- 
tion. It is needed not only as an aid to 
making the final choice of lixiviant, but 
also for evaluating well construction and 
completion techniques and for demonstrat- 
ing restoration procedures. 

Types of Tests 

Pilot-scale tests can be divided 
into two classifications. The first type 
is called push-pull, or huff-and-puff. 
The lixiviant is injected and recovered 
from the same well. The second type can 
be called flow-through. The lixiviant is 
injected, flows through the formation, 
and is recovered from other wells. 

Consultants disagree as to the value 
of push-pull tests. Some believe that 
these tests provide a reliable and rela- 
tively economical method of evaluating 
the lixiviant under field conditions. 
Professor Robert Schechter of the 
University of Texas at Austin determined 
that one well serving first for injection 
and later for production, together with 
two observation wells, would provide 



18 



sufficient data for a field evaluation of 
the chloride preflush previously des- 
cribed. The well pattern is an L, with 
the injection-production well at the 
corner. 

A push-pull test can be misleading, 
especially one with no observation wells, 
if adsorption is not considered. 
Adsorbed species do not move as far from 
the well as would be calculated if 
fluid volume were considered. In 
field test, most of the ammonium 
remained within one- fourth of the 
contacted by the fluid {6). In 
case, the number of pore volumes 



only 

one 

ions 

ore 

that 

of 

groundwater sweeping required for resto- 
ration of groundwater quality would have 
been underestimated by a factor of 4 if 
adsorption had been neglected. 

Flow- through tests resemble most 
commercial operations more closely than 



do push-pull tests, 
tion may be forced 
method of leaching 
lenses and pinchouts 
logic communication 
the flow- through method 
far more common. 



A commercial opera- 
to use a push-pull 

if numerous clay 

prevent good hydro- 

between wells, but 

is preferred and 



Problems and Solutions 

Because problems can occur that ren- 
der a pilot field test useless as a guide 
in making the final choice of lixiviant, 
these problems are briefly described 
below. Problems that have occurred in 
past tests include the following: 

1. Leaking casings. 

2. Clogging of well screens or 
nearby formation. 

3. Clogging of formation near a 
production well . 

4. Reprecipitation of uranium. 

The most serious leaks are caused by 
well completion tools gouging holes in 
PVC casing. For example, leaks have been 
caused by underreamers when the tool has 
been pushed down or pulled up with the 
blade not fully retracted. Sand can 
interfere with the retracting of the 
blade. Damage can also be caused when 
any tool that fits tightly is forced past 
casing that is somewhat curved because of 
hole deviation. 



Most operators believe that a flow- 
through pilot test before beginning a 
flow-through commercial operation is 
advisable even if a push-pull test has 
been used for a preliminary evaluation. 
A common configuration is a five-spot 
pattern, injecting at the corners and 
producing at the center, because usually 
there is more fluid resistance to injec- 
tion than to production. The results of 
flow-through tests can be quite different 
from those of push-pull tests. In unpub- 
lished flow-through tests at one site, 
increases in uranium concentration and Eh 
lagged behind the changes in the other 
components. Over three pore volumes of 
lixiviant were injected before the ura- 
nium concentration and Eh rose. However, 
in a previous push-pull test at the same 
site, that lag was not observed. There- 
fore, the push-pull test resulted in an 
overly optimistic prediction about the 
ease of leaching. 



Smal 
penetrate 
where a 
improper 
immobile 
but the 
the way 
has publi 
ing casin 



ler leaks can 

all the way 

joint is 1 

gluing. Hoi 

with screws 

screws should 

through the 

shed other i 

gs (1). 



occur where screws 
through a joint, or 
eaking because of 
ding a glued joint 
may be advisable, 
not penetrate all 
joint. The Bureau 
nformation concern- 



Clogging of well screens or the 
nearby formation can occur before lixivi- 
ant is injected as a result of wall cake 
left from drilling. Techniques that pro- 
duce acceptable exploration holes do not 
necessarily produce acceptable injection 
wells. Care should be taken to avoid 
drilling fluids that leave a thick wall 
cake. If underreaming or perforating 
is used for well completion, then the 
problems from the well cake are mini- 
mized. However, if conventional well 
screens are used, then polymer drilling 



19 



fluids or bentom'te with a polymer 
additive generally allow better injectiv- 
ity than just bentonite. The Bureau has 
prepared more detailed publications on 
this topic a, 33-34). 

Clogging of the formation near an 
injection well screen can occur upon the 
injection of lixiviant incompatible with 
the formation or groundwater. An example 
is clay swelling caused by injecting 
sodium carbonate-bicarbonate into a for- 
mation that has a high concentration of 
swelling clay and where the preleach 
groundwater is low in sodium. Usually, 
the incompatibility can be identified in 
laboratory column leaching tests. 

Clogging of the formation near a 
production well can be caused by precipi- 
tation of calcium carbonate. The pres- 
sure near a production well is lower than 
in the rest of the formation, so some of 
the carbon dioxide can come out of satu- 
rated solution. Apart from the gas phase 
blockage, the lowering of dissolved car- 
bon dioxide can contribute to clogging by 
raising the pH, thereby decreasing the 
solubility of calcium carbonate and per- 
haps causing its precipitation. The 
chloride preflush appears to be a promis- 
ing method of minimizing this problem. 
An alternative approach is to begin 
injection with a high pH lixiviant, to 
precipitate calcium in place so it cannot 
become concentrated near the production 
well. This approach has the disadvantage 
that uranium will be coprecipitated (24) 
and become less accessible to tFe 
lixiviant. 

Clogging of production well screens 
or the nearby formation can also occur as 
a result of fines migration. Periodic 
vigorous well development such as air- 
lifting may be required. 

Reprecipitation of uranium can occur 
if the strength of the lixiviant (pH, Eh, 
or concentration) changes excessively as 
the lixiviant moves from injection to 
production well. This problem can be 
minimized by increasing the strength of 
the lixiviant gradually as leaching 
begins and by selecting a lixiviant com- 



patible with the formation. The problem 
can be reduced by reducing well spacing., 
but that is an expensive solution. 

GEOCHEMICAL MODELS 

The geochemistry associated with in 
situ leaching of uranium is extremely 
complex and is difficult to characterize 
without the use of a computer. Geochemi- 
cal computer models can be divided into 
two major categories. The first type of 
model, the equilibrium approach, is use- 
ful for describing numerous interactions 
of a complex system of aqueous species 
and solid phases. Equilibrium programs 
generally require a complete chemical 
analysis of a solution including field 
measurements of Eh, pH, and temperatureo 
The output of this type of model can be 
used to determine the reactions that are 
likely to occur within a given system but 
gives no information concerning the rates 
of the reactions. Therefore, equilibrium 
models cannot adequately describe time- 
dependent reactions, which are affected 
by fluid velocity and dispersion. 

The second typ9 of model , the 
kinetic model, simulates the progress of 
kinetic reactions as a function of time 
and location (2^-3^, 26). Kinetic models 
are coupled with hydrology models and 
thus account for formation factors such 
as porosity, permeability, and dispersi^f- 
ity. A hydrology model adaptable to such 
coupling has been developed by Schmidt 
(27). The main disadvantage of kinetic 
models is that they are limited to rela- 
tively few chemical reactions. 

In addition to hydrologic data for 
the aquifer, kinetic models require a 
large amount of chemical information per- 
tinent to the reactions being modeled. 
Chemical data include reaction rate coef- 
ficients, which are determined by exten- 
sive laboratory experiments. Once the 
pertinent physical and chemical proper- 
ties have been established, an operator 
can determine the effect of well 
pattern and spacing, pumping and 
injection rate, and injected oxygen con- 
centration on uranium production. 
Because kinetic models cannot be used in 



20 



selecting a suitable lixiviant and 
oxidant unless pertinent reaction rates 
are first determined through laboratory 
experiments, this report will concentrate 
on equilibrium modeling. 

Numerous equilibrium models have 
been developed in recent years for the 
purpose of modeling geochemistry of natu- 
ral water. A useful summary and compari- 
son of these models has been published by 
Nordstrom (22). Although these models 
can be applied to a wide range of hydro- 
geochemical problems, most of them are 
unable to model reactions associated with 
in situ leaching of uranium roll -front 
deposits. As of 1980, an updated version 
of WATEQF (23, 25^) is the most suitable 
program for~l[escribing the aqueous geo- 
chemistry of uranium. 

The original version of this program 
(WATEQ) was written in PL/1 and was pub- 
lished by Truesdell and Jones (31) in 
1973. In 1976, Plummer (23) publTshed a 
revised version, WATEQF, in the program- 
ming language FORTRAN IV. This revised 
version recently has been enlarged by 
Runnel Is (25^) to include many dissolved 
species and solid compounds of uranium, 
vanadium, and molybdenum. 

Detailed explanations of theory and 
operation for WATEQ and WATEQF are pro- 
vided by Truesdell and Jones (31); 
Plummer (23); and Lueck (17^). Briefly, 
the purpose of WATEQF is to solve a large 
number of simultaneous equations that 
describe all of the known equilibrium 
reactions that may occur in a given solu- 
tion. The program requires as input a 
relatively complete chemical analysis of 
the solution of interest. This analysis 
includes temperature. Eh, pH, concentra- 
tions of all major cations and anions, 
and concentrations of the minor and trace 
elements that are relevant to the inves- 
tigation. These data are summarized on 
the first page of output (table A-1). In 
addition to the chemical analysis of the 
solution, a table of thermodynamic data 
for all reactions modeled by the program 
must also be read into the computer. 
This feature enables the user to update 



thermodynamic data without having to 
modify the program itself. 

After the data are read, the program 
adjusts the equilibrium constants for 
temperature effects using the Van't Hoff 
equation or a user defined analytical 
expression. It then uses an iterative 
procedure to determine the activities and 
concentrations of all aqueous species. 
At the start of each iteration, activity 
coefficients are calculated from the 
Davies equation, the extended Debye- 
Huckel equation, or the Guntelberg equa- 
tion. Then a back substitution procedure 
is used to solve mass action and mass 
balance equations for dissolved species. 
When the computed concentrations of major 
anions agree with the respective analyti- 
cal values within a specified tolerance, 
iteration stops and the final activities 
are retained for computing the stabili- 
ties of solid phases. 

The second page of output 
(table A-2) lists the differences between 
computed and analytical values of major 
anions for each iteration. It also 
describes the solution in terms of 
important parameters such as ionic 
strength, total dissolved solids, pH, Eh, 
electron activity, and partial pressures 
of dissolved gases, {O2, CO2, and CH4). 
The concentrations, activity coeffi- 
cients, and activities of aqueous species 
are tabulated on pages 3-6 of the output 
(table A-3). This distribution of spe- 
cies list yields valuable information on 
the speciation of dissolved constituents. 
For example, the sample output indicates 
that about 59 pet of the total molality 
of dissolved calcium is free ionic Ca"*""*"; 
about 40 pet is aqueous CaS04; and about 
1 pet is CaHC03. 

After the distribution of aqueous 
species is determined, WATEQF computes 
the state of saturation of the solution 
with respect to various minerals and 
amorphous solid compounds. The diagnos- 
tics that describe the stabilities of the 
solid phases constitute the final section 
of output (table A-4). When interpreting 
these diagnostics, it is important to 



21 



remember that each phase in the first 
column actually represents a reaction 
involving that phase with the appropriate 
aqueous species. These reactions are 
included in several references describing 
WATEQ or WATEQF (r7, 23, 25, 31^). 

To determine the stability of a 
solid phase within a solution, it is nec- 
essary to compare the activity product of 
the ions involved in the appropriate 
reaction (ion activity product, or 
i.a.p.) to the thermodynamic equilibrium 
constant (K^) for that reaction. This 
comparison can be expressed as a ratio 
(i.a.p./K-,.) or as a log of that ratio 
(log 10 (i.a.p./K^)). The log of the 
ratio is termed the saturation index 
(S.I.) and is a useful indicator of a 
mineral's stability in a solution. 
Another useful indicator is the Gibb's 
free energy of reaction (AGp), which is 
the amount of energy (expressed in kilo- 
calories) that must be supplied to allow 
the reaction to proceed. The following 
guidelines should be used in interpreting 
saturation indices and free energies of 
reaction: 

1) S.I. and AGp less than zero: 
The reaction will proceed spontaneously, 
although the rate may be extremely slow. 
Since the reactions modeled by this pro- 
gram describe the dissolution of solid 
phases, this would indicate that the 
solution is undersaturated with respect 
to the solid and that the solid (if pres- 
ent) would go into solution. A highly 
negative S.I. or aG^ may indicate that a 
particular mineral does not exist in the 
system. 

2) S.I. and aG^- close to zero: The 
reaction is at equilibrium. 

3) S.I. and AGr- greater than zero: 
The safest interpretation of this condi- 
tion is that the solution is supersatu- 
rated with respect to a given mineral and 
that the mineral is stable in that par- 
ticular aqueous environment. This does 
not indicate precipitation. The formal 
interpretation of a positive aG^- is that 
the reaction cannot proceed unless energy 
is supplied from an external source. In 



some cases, positive values suggest that 
a particular solid may precipitate from 
solution, but caution should be used in 
this type of interpretation. It is pos- 
sible for solutions to remain supersatu- 
rated with respect to some minerals for 
long periods of time without precipita- 
tion. Some minerals are not known to 
precipitate from solution at all. 

The modified version of WATEQF has 
had many applications related to the gen- 
esis of uranium and vanadium deposits, 
hydrogeochemical exploration for uranium, 
and in situ leaching uranium roll front 
deposits (17, 25^). In every application^ 
it allowed~Tnterpretations that could not 
have been made from chemical data aloneo 
Application of WATEQF to in situ leaching 
of uranium can provide operators with 
valuable information on the speciation of 
uranium and vanadium complexes in solu- 
tion; the solubility of uranium and vana- 
dium minerals; the formation of gaseous 
oxygen and carbon dioxide, which could 
reduce permeability, and the precipita- 
tion of solid phases, which could reduce 
permeability or remove uranium from 
solution. 

WATEQF is more useful for predicting 
changes in solubility than for predicting 
the solubility itself. This is because 
supersaturations of 10 to 100 are often 
found, so the calculated solubilities 
are not reliable predictors of 
concentrations. 

As of 1980, at least two companies 
are using WATEQF (as modified by Runnels) 
to assist in determining how the lixivi- 
ant composition should be changed to 
improve leaching. One company uses it to 
help select the most cost-effective lix- 
iviant composition for dissolving the 
uranium minerals. The cost of a solution 
providing a given pH and Eh can be esti- 
mated, and the solubilities of the miner- 
als can be predicted with WATEQF. Thus, 
for a given lixiviant cost, the program 
can help select the combination of pH and 
Eh maximizing solubility. Judgment is 
still required for balancing cost versus 
solubility, however. 



22 



Another company uses WATEQF to pre- 
dict whether solubilities will increase 
or decrease with changes in carbonate 
concentration, pH, or Eh. It is espe- 
cially helpful in determining the 
probable cause and suggesting a cure when 
pilot tests are yielding much less ura- 
nium than expected. This company also 
uses WATEQF to predict the relative 
amounts of uranium species. Uranium as a 
monocarbonate complex will not load on 
anionic exchange resins and so is unde- 
sirable. WATEQF predicts what fraction 
will be in monocarbonate, di carbonate, 
and tri carbonate complexes. The program 
has also been used to predict fouling 
from minerals precipitating in pipes and 
to study restoration geochemistry. 

An additional application of WATEQF 
can be demonstrated through the following 
example supplied by Professor Donald 
Runnels of the University of Colorado 
(25). At the 1979 Symposium on the 
Grants Mineral Belt, chemical evidence 
was presented in support of the existence 
of a calcium-uranium-phosphate associa- 
tion in uranium ores (personal communica- 
tion with Dr. Runnels, Sept. 22, 1980). 
Results of geochemical modeling suggest a 
potentially important role of dissolved 
phosphate in controlling the solubility 
of this association in acid lixiviants 
and in affecting the recovery of uranium 
from in situ operations. 

The data for this example consist of 
complete chemical analyses of solutions 
that were collected from observation 
wells at the Rocky Mountain Energy Co.'s 
Nine Mile Lake Site, Wyoming. Although 
the concentration of dissolved phosphate 
never exceeds 13 ppm, many of the solu- 
tions appear to be supersaturated with 
respect to the mineral ningyoite, 
(U,Ca)2(P04)2 • 1 - 2H2O (20). Ningyoite 
and other solid compounds of uranium and 
phosphate appear to be insoluble under 
most conditions of the sulfuric acid lix- 
i VI ant, except at very high Eh. This 
suggests that an operator might be able 
to increase the efficiency of the leach- 
ing process by reducing the concentration 
of phosphate in the lixiviant. 



At the present time, WATEQF seems to 
be the most suitable equilibrium model 
for problems related to in situ leaching 
of uranium because it includes reactions 
for minerals associated with uranium 
deposits. However, it is limited by its 
inability to model mass transfer, adsorp- 
tion, or solid-solution reactions. Other 
models having one or more of these capa- 
bilities include REDEQL2 (18) and EQ3/6 
(36). These models, however, must be 
mo3"ified to include reactions associated 
with in situ leaching of uranium. 

SUMMARY 

The selection of a lixiviant pro- 
ceeds through three phases. First, gen- 
eral advantages and disadvantages of 
lixiviants are considered. These general 
considerations include technical, eco- 
nomic, and environmental factors. Cur- 
rently, restoration of groundwater 
quality are causing a movement away from 
ammonium carbonate-bicarbonate toward 
sodium bicarbonate and dissolved carbon 
dioxide. The cost of the oxidizer should 
be carefully considered, because it can 
exceed the cost of all the other 
chemicals. 

Second, lixiviants that seem promis- 
ing are tested with ore (cores) from the 
site to be leached. Laboratory batch and 
column leaching experiments measure 
leaching efficiency, consumption, and 
effect on permeability. These tests can 
be misleading if not conducted and inter- 
preted with care. 

Third, a pilot-scale field test is 
conducted. Proper well construction is 
vital to the success of this test. The 
test can be either the push-pull or flow- 
through type. The former is cheaper, but 
the latter simulates most commercial 
operations more closely. 

Computer modeling of the geochem- 
istry can aid in the selection process. 
Such models are being used by at least 
two leaching companies to predict changes 
in solubilities associated with possible 
changes in lixiviant composition. 



23 



REFERENCES^ 



J. 



K.,.NI. 



T. Nigbor, and 



Ahlness. 

D. R. Tweeton. Drilling Fluids and 
Well Casing Materials for In Situ 
Uranium Leaching. Pres. at the 
Fall Meeting, Soc. Mining Eng., 
AIME, Tucson, Ariz., Oct. 17-19, 
1979, 26 pp.; available upon 
request from J. K. Ahlness, Bureau 
of Mines, Minneapolis, Minn. 

Bommer, P. M. A Streamline- 
Concentration Balance Model for In 
Situ Uranium Leaching and Site 
Restoration. Ph.D. Dissertation, 
Univ. of Texas, Austin, 1979, 
262 pp. 

Bommer, P. M., and R. S. Schechter. 
Mathematical Modeling of In-Situ 
Uranium Leaching. Pres. at 53d 
Ann. Meeting, Soc. Petrol. Eng., 
AIME, Houston, Tex., Oct. 1-3, 
1978, SPE Preprint 7533, 6 pp. 

Chemical Marketing Reporter, Schnell 
Pub. Co., Inc., Sept. 22, 1980, 
43 pp. 

Davy, R. The Goechemical and Radio- 
metric Behavior of Uranium and 
Other Radioactive Elements in Rela- 
tion to Exploration for Secondary 
Uranium. Amdel Bull., v. 17, April 
1974, p. 26. 

Drever, J. I., and C. R. McKee. The 
Push-Pull Test. A Method of Evalu- 
ating Formation Adsorption Param- 
eters for Predicting the Environ- 
mental Effects on In-Situ Coal Gas- 
ification and Uranium Recovery. In 



Situ, V. 4, No. 3, 1980, pp. 196- 
199. 

Engelmann, W. H., P. E. Phillips, 
D. R. Tweeton, K. W. Loest, and M. 
T. Nigbor. Restoration of Ground- 
water Quality Following Pilot- 
Scale Acidic In Situ Uranium 
Leaching at Nine-Mile Lake Site 
Near Casper, Wyoming. Pres. at 
55th Ann. Meeting, Soc. Petrol. 
Eng., AIME, Dallas, Tex., 
Sept. 21-24, 1980, SPE Preprint 
9494, 24 pp. 



FMC Corp., Industrial 
Hydrogen Peroxide, 
dated), p. 3. 



Chemical Div. 
New York, (not 



9. FMC Corp., Industrial Chemical 
Group. Sodium Sesquicarbonate. 
Philadelphia, Pa., Technical data 
sheet Dl-18, 1977, 1 p. 

10. Grant, D. C. In Situ Leaching 

Studies of Uranium Ores--Phase IV. 
BuMines Open File Rept. 52-79, 
1978, 497 pp.; available from 
National Technical Information 
Service, Springfield, Va., 
PB 296336/AS. 

11. Hartman, G. J., and G. J. Catchpole. 

Groundwater Restoration of the 
OPI-Western Joint Venture Orebody. 
In Uranium Resource Technology, 
Seminar III, Colorado Sch. Mines, 
Golden, Colo., June 1980, 
pp. 173-185. 

12. Jacobucci, J. R. Private Communica- 

tion to D. R. Tweeton, Bureau of 
Mines, Minneapolis, Minn. 



are ava i I ab I e 
the Bureau of 
Albany, Oreg,, 



^References 10 and 13 
for consultation at 
Mines libraries in 
Avondale, Md . , Denver, Colo,, Pitts- 
burgh, Pa,, Reno, Nev , , Ro I I a , Mo,, 
Salt Lake City, Utah, Spokane, Wash,, 
Tuscaloosa, Ala,, Twin Cities (Minnea- 
polis), Minn,, Boulder City, Nev,, and 
at the National Library of Natural 
Resources, U,S, Department of the 
Interior, Washington, D,C, 



13, Kasper, D, R,, 
L, D, Munsey, 
C. K, Chase. 
Assessment of 
BuMines Open 
1979, 292 pp. 
National 
Service, 



H. W. Martin, 

R, B, Bhappu, and 

Environmental 

In Situ Mining. 

File Rept. 101-80, 

; available from 

Technical Information 

Springfield, Va., 



PB 81 106783. 



24 



14. Kirkham, 



R. M. 



Promises and 22. 



15. 



16. 



17. 



18. 



19. 



20. 



21. 



Problems of a "New" Uranium Mining 
Method: In Situ Solution Mining. 
Colorado Geological Survey (Den- 
ver, Colo.), Environmental Geol- 
ogy II, 1979, 21 pp. 



Larson, W. 
Leach Minin 
BuMines IC 

Litz, L. M. 
With Oxygen 
Ann. Meeti 
Feb. 14-18, 
8 pp. (Th 
published 
corrected 
communicati 



C. Uranium In Situ 
g in the United States. 
8777, 68 pp. 

In-Situ Uranium Mining 

Pres. at 1980 AIME 

ng. Las Vegas, Nev., 

1980, Preprint 80-71, 

e formula as originally 

contained an error, 

through private 

on.) 



Needes, C. R. S., M. J. Nicol, and 
N. P. Finkel stein. Electrochemi- 
cal model for the leaching of Ura- 
nium Dioxide: 2-Alkaline Carbon- 
ate Media. Ch. in Leaching and 
Reduction in Hydrometallurgy, ed. 
by A. R. Burkin. IMM, London, 
1975, pp. 12-19. 



23. 



24. 



25. 



Lueck, S. L. Computer Modeling of 
Uranium Species in Natural Waters. 
M. S. Thesis, Dept. Geol . Sci., 
Univ. Colorado, Boulder, 1978, 
170 pp. 

McDuff, R. E., and F. M. Morel. 
Description and Use of the Chemi- 
cal Equilibrium Program REDEQL2. 
W. M. Keck Lab. of Environmental 
Eng., Tech. Rept. EQ-73-02, Calif. 
Inst. Techno!., Pasadena, Calif., 
1975, 83 pp. 

Merritt, R. C. The Extractive 
Metallurgy of Uranium. Colorado 
Sch. Mines Res. Inst., Golden, 
Colo., 1971, 576 pp. 



Muto, T., R. Meyrowitz, A. M. Pom- 
mer, and T. Murano. Ning- 26, 
yoite, a New Uranous Phosphate 
Mineral From Japan. Am. Mineral- 
ogist, V. 44, May-June 1959, 
pp. 663-650. 



27. 



Nordstrom, D. K., and others. A 
Comparison of Computerized Chemi- 
cal Models for Equilibrium Calcu- 
lations in Aqueous Systems. Ch. 
in Chemical Modeling in Aqueous 
Systems. Speciation, Sorption, 
Solubility, and Kinetics--Am. 
Chem. Soc. Symp. Series, 1979, 
pp. 857-892. 

Plummer, L. N., B. F. Jones, and 
A. H. Truesdell. WATEQF— A 
FORTRAN IV Version of WATEQ, a 
Computer Program for Calculating 
Chemical Equilibrium of Natural 
Waters. U.S. Geol. Survey Water 
Res. Inv. 76-13, 1976, 61 pp. 



Potter, R. W., 


M. 


A. Clynne, 


J. M. Thompson, 


V. 


L. Thurmond, 


R. C. Erd, 


N. 


L. Nehring, 


K. A. Smith, 


P. 


J. Lamothe, 


J. L. Seeley, 


D. 


R. Twee ton. 


G. R. Anderson, 


and 


W. H. Engel- 



mann. Chemical Monitoring of the 
In-Situ Leaching of a South Texas 
Uranium Orebody. U.S. Geol. Sur- 
vey Open File Rept. 79-1144, 1979, 
p. 25. 



Runnells, D. D., 
S. L. Lueck, 
Applications of 
to the Genesis, 
In-Situ Mining 
Vanadium Deposits 
Bur. Mines and 



R. Lindberg, 

and G. Markos. 

Computer Modeling 

Exploration, and 

of Uranium and 

New Mexico 

Miner. Res. 



Mem. 38, 
20 pp. 



Socorro, N. Mex., 1981, 



Schechter, R. S., and P. M. Bommer. 
Optimization of Uranium Leach Min- 
ing. Pres. at 55th Ann. Meeting, 
Soc. Petrol. Eng., AIME, Dallas, 
Tex., Sept. 21-24, 1980, SPE Pre- 
print 9494, 6 pp. 

Schmidt, R. D. Computer Modeling of 
Fluid Flow During Production and 
Environmental Restoration Phases 
of In Situ Uranium Leaching. 
BuMines RI 8479, 1980, 43 pp. 



25 



28. Stephen, H., and T. Stephen. Solu- 

bilities of Inorganic and Organic 
Compounds. V. 1, Binary Systems, 
Part 1. The Macmillan Co., New 
York, 1963, 960 pp. 

29. Thiede, D. M., and D. W. Walker. 

South Texas Uranium Leach Drilling 
and Completion Technology. Pres. 
at South Texas Minerals Section of 
AIME Uranium In Situ Symposium, 
Corpus Christi, Tex., September 
1977, 12 pp.; available upon 
request from D. R. Tweeton, Bureau 
of Mines, Minneapolis, Minn. 

30. Thompson, W. E., W. V. Swarzenski, 

D. L. Warner, G. E. Rouse, 
0. F. Carrington, and R. Z. Pyrih. 
Ground-water Elements of In Situ 
Leach Mining of Uranium. Prepared 
for U.S. Nuclear Regulatory Com- 
mission, August 1978, 173 pp.; 
available from National Technical 
Information Service, Springfield, 
Va., NUREG/CR-0311. 

31. Truesdell, A. H., and B. F. Jones. 

WATEQ, a Computer Program for Cal- 
culating Chemical Equilibria of 
Natural Waters. U.S. Geol . Sur- 
vey, 1973, 73 pp.; available from 
National Technical Information 
Service, Springfield, Va., 
PB 220 464. 



34. Tweeton, D. R., and K. Connor. Well 

Construction Information for In 
Situ Uranium Leaching. BuMines 
IC 8769, 1978, 19 pp. 

35. Tweeton, D. R., T. Gui linger, 

W. M. Breland, and R. S. Schech- 
ter. The Advantages of Condition- 
ing an Orebody With a Chloride 
Solution Before In-Situ Uranium 
Leaching With a Carbonate Solu- 
tion. Pres. at 55th Ann. Meeting, 
Soc. Petrol. Eng., AIME, Dallas, 
Tex., Sept. 21-24, 1980, SPE Pre- 
print 9490, 7 pp. 

36. Wol ery, T. J. Calculation of Chemi- 

cal Equilibrium Between Aqueous 
Solution and Minerals: The EQ3/6 
Software Package. UCRL-52658, 
Lawrence Livermore Lab., Univ. 
California, 1979, 41 pp. 



32. Tsoung, Y. Calcite Control in 

In-Situ Leaching Process. Pres. 
at 54th Ann. Meeting, Soc. Petrol. 
Eng., AIME, Las Vegas, Nev., 
Sept. 23-26, 1979, SPE Pre- 
print 8319, 7 pp. 

33. Tweeton, D. R., G. R. Anderson, and 

W. H. Engelmann. Bureau of Mines 
Research in Injection Well Con- 
struction and Environmental 
Aspects of In Situ Uranium Leach- 
ing. Pres. at 1978 AIME ann. 
Meeting, Denver, Colo., Feb. 26- 
Mar. 2, 1978, Preprint 78-AS-lll, 
8 pp. 



26 



APPENDIX 

All pages in the appendix are printouts by the geochemical modeling program 
WATEQF. As indicated in "Geochemical Models," they provide examples of required 
input data and of the resulting output. 



27 



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